Author: Arnie Benn

RETURN to Periodic Table
Below: Electron Shell, Bonding & Ion Formation, Magnetic Properties

Scandium is the 21^st element on the periodic table. It has 21 protons and 24 neutrons for a mass of 45 amu, and 21 electrons.

Electron Shell

DISCLAIMER: This electron domain geometry is designed to show the most symmetrical resonance structure for the orbitals, according to the Quicycle model. This is not the only resonance structure possible for this element. (In this series we are also not differentiating between high-spin and low-spin states. These variations should be extrapolated from the geometries shown.)

Scandium is the first element to feature electrons in the d–orbital. Since d-electrons are resonating in the same space as p-electrons within their parent s-shell, we propose that these orbitals are able to hybridize together in search of lowest energy and highest stability. [Ref] If that is the case, scandium would feature a 3^rd shell containing 3 di-electrons and 1 unpaired electron in p³d-hybridized tetrahedral symmetry. This pd-hybridization does not need to involve the 3s-orbital electrons, and they can return to their preferred spherical di-electron state. (see image below)

The 3^rd shell tetrahedral symmetry will also align antiparallel (and antiprismatic) to the 2^nd shell tetrahedral electron arrangement within, in order to minimize di-electron repulsion between shells.

NOTE: The small spheres in the image above simply indicate the directions of maximum electron density. The 3^rd shell hybrid orbitals themselves will assume a spherical tetrahedral structure that divides the shell into four roughly equal volumes. Each shell segment will be filled with electron density. It will be highest at the center of the face of each orbital (as in the traditional sp³ lobe shapes) and will decrease toward the nodal regions between orbitals — as wave structures usually do — where electron density will be lowest (though not zero).

NOTE: Even though it is often useful to talk about these orbitals as separate, they are all — the entire atom is — part of a single, coherent, harmonic, resonant, phase-locked, spherically-symmetrical quantum wave state, and it is all electromagnetic at the root-energy level. Orbitals and their ‘boundaries’ can be seen as nothing more than nodes and antinodes in this harmonic wave structure.

NOTE: The model’s claim is not that an isolated atom always has a definite antinode location in the lab frame. It is that, conditional on one electron being found in a given orbital ‘antinode’ region or direction, the other electrons’ positions must correlated with it in a specific geometric pattern. This conditional-correlation structure is a body-frame feature: it describes the relative positions of electrons within the atom, not their absolute positions in the lab frame. It is the internal frame of reference in which the shell’s wave resonance condition closes and its standing-wave pattern is organized. For a free atom in an unprepared ensemble, all orientations are statistically equivalent so the observed density is typically an orientation average over body-frame orientations, which is observed in the lab frame to be spherically symmetrical.

An alternative representation of the scandium atom is given below. It shows a 3^rd shell p³d-hybridised tetrahedral symmetry that is similar to the electron configuration of chlorine — 3 di-electrons and 1 unpaired electron.

A key difference, though, is that these hybrid orbitals exist within the added sheath of the 3s²-orbital di-electron, since it is not needed for hybridization. It is proposed that the 3s² di-electron gives additional stability to the tetrahedral p³d-hybrid orbitals superimposed within it — like harmonic frequencies upon a fundamental — even though it contains an unpaired electron.

The remaining elements in the d-block will also experience this 3s-orbital stabilization phenomenon, since they also achieve at least 4-directional symmetry with only their p– and d-orbitals.

BACK TO TOP

Bonding & Ion Formation

The 4s²-orbital di-electron shell completes the isolated neutral atom. When removing two electrons, the free-ion Sc²⁺ is formed. When forming a metallic crystal, some of the 4s and 3d electron density delocalizes to form the metallic bond. In both cases, the core electron domain geometry remains the same. (A third ionization, however, would change it, possibly into an sp³-hybridized tetrahedral structure.)

(In a few cases, such as chromium (Cr) and copper (Cu), the 4s-orbital outer shell contains only 1 electron, since the other joins the 3rd shell hybridization within in order to help it achieve 8-directional symmetry.)

BACK TO TOP

Magnetic Properties

Scandium has one unpaired electron. In all previous elements, unpaired electrons have occurred in the outer (valence) shell of the atom. Those are the electrons involved in chemical reactions. Scandium is the first case of an atom with an unpaired ‘core’ valence electron. It is protected from reacting (at least, until scandium becomes an Sc²⁺ ion), and it is stabilized by the 3s di-electron ‘fundamental.’ We propose that it is this protection and stabilization of unpaired valence electrons on the core that allows them to interact magnetically, and that consequently determines an element’s magnetic properties.

UNPAIRED ‘CORE’ VALENCE ELECTRONS:
While these electrons are technically still valence electrons, which contribute electron density to the crystal electronic band structure in the conduction electron matrix — the 3D electron gas — of the solid metal crystal, some of the 3d electron density remains concentrated around the atomic cores.

PARAMAGNETISM:
In the presence of an external magnetic field, an unpaired core electron will orient its spin to align with the magnetic field. This will cause the atom to be drawn into and toward that field — via magnetic field cancellation — giving it a positive magnetic susceptibility (χ_m) value. This attractive force is called paramagnetism, and its effects only last as long as the external magnetic field is present. We might therefore presume that, the stronger the paramagnetism, the more ‘unpaired electron character’ is present. Surprisingly, this does not seem to go according the number of unpaired electrons present, as the diagram below illustrates. There must therefore be other contributing factors, which we will investigate below.

Despite only have a single unpaired electron, scandium is the second most paramagnetic metal of the 3d row — excluding the ferromagnetic iron, cobalt, and nickel.

As proposed above, scandium has one unpaired electron in one position of a tetrahedral electron structure. The images below are intended to represent the orientation of this 3^rd shell electron in an external (or adjacent atom’s) magnetic field (whose north pole is pointing upwards). (The purple arrow inside the atom represent the unpaired electron and the black dots represent the relative positions of the 3 di-electrons in that shell.)

It is here proposed that scandium represents a rather simple case of one electron that is susceptible to an external magnetic field, and whose effects are not mitigated by other unpaired electrons within the core electron topology. As we will see in the case of some other d-block metals, the presence of unpaired electrons spin-bonding within an atomic shell can cancel magnetic field and reduce an electron’s effective magnetic signature with respect to an external magnetic field.

PARAMAGNETIC STRENGTH ANALYSIS:
The following diagram shows the relative paramagnetic strengths of the transition metals, along with their proposed hybrid orbital geometries. (See paramagnetic strength trend analysis for more detail.)

---

OTHER PARAMAGNETIC 3d METALS: Scandium, Titanium, Vanadium, Chromium (also antiferromagnetic), Manganese
FERROMAGNETIC 3d METALS: Iron, Cobalt, Nickel
DIAMAGNETIC 3d METALS: Copper, Zinc

RETURN to the Periodic Table

BACK TO TOP

RETURN to Periodic Table

Calcium is the 20^th element on the periodic table. It has 20 protons and 20 neutrons in the nucleus, giving it a mass of 40 amu, and it has 20 electrons enveloping the nucleus.

Electron Shell

Calcium has the same electron configuration as magnesium, but with three full shells within that have the identical configuration to argon. Being one shell larger than magnesium, calcium’s valence electrons are less well bound, causing it to be more reactive than magnesium, but much less so than potassium.

As we saw in the case of argon, the 3^rd shell orbitals are more like spherical tetrahedra, and the 4^th shell is a di-electron in a spherical s-orbital. These orbitals represent phase-locked, resonant, coherent, harmonic, stationary waves.

Ion formation

Calcium will give up its valence electrons in an ionic interaction in order to reach the stability of the 3s²3p⁶ noble gas configuration of argon, which is a multi-di-electron state with three concentric full shells. That is why calcium forms a 2+ ionic state.

---

RETURN to the Periodic Table

OTHER GROUP II ELEMENTS: Beryllium, Magnesium, Calcium, Strontium, Barium

RETURN to Periodic Table

Potassium is the 19^th element on the periodic table. It has 19 protons and 20 neutrons in the nucleus, giving it a mass of 39 amu, and it has 19 electrons enveloping the nucleus.

Electron Shell

Potassium has the same electron configuration as sodium, but with three full shells within that have the identical configuration to argon. Being one shell larger than sodium, potassium has a lower ionization energy and is therefore more reactive.

As we saw in the case of argon, the 3^rd shell orbitals are more like spherical tetrahedra, and the 4^th shell is a single electron in a spherical s-orbital. These orbitals represent phase-locked, resonant, coherent, harmonic, stationary waves.

Ion formation

With a lower ionization energy, potassium will give up its valence electron more eagerly than sodium in an ionic interaction, in order to reach the stability of the 3s²3p⁶ noble gas configuration of argon, which is a multi-di-electron state with three concentric full shells. That is why potassium forms a 1+ ionic state so violently.

Pure potassium metal reacts explosively when placed in water as it donates its valence electron to oxygen. The heat of this reaction ignites the hydrogen gas that is also produced from the water. Potassium burns with a purple flame and throws purple sparks. It is used to create this same effect and color in fireworks. A fun video showing potassium metal and its reactivity can be found HERE.

---

RETURN to the Periodic Table

OTHER GROUP I ELEMENTS: Lithium, Sodium, Potassium, Rubidium, Cesium

RETURN to Periodic Table

Argon is the 18^th element on the periodic table. It has 18 protons and 22 neutrons in the nucleus, giving it a mass of 40 amu, and it has a full shell of 18 electrons enveloping the nucleus. Its full electron shell causes it not to react with other atoms. It is a gas under normal conditions, and it is called a ‘noble’ gas because of its non-interaction.

Electron Shell

Argon has three full electron shells. With 6 electrons in a 3p orbital, argon is believed to achieve stability with octahedral symmetry, and its 3s electrons can unhybridize and return to their preferred spherical di-electron state. It is proposed in the theory of Sub-Quantum Chemistry, however, that argon is more stable with full shells containing 4 di-electrons in sp³-hybridization (shown below).

This high degree of symmetry, with all electrons paired, renders argon, like helium and neon, unreactive.

NOTE: The small spheres in the image above simply indicate the directions of maximum electron density. The 3^rd shell orbitals themselves are spherical tetrahedra, not spheres. Only s-orbitals are spherical. It is noteworthy that the tetrahedral arrangements in adjacent shells align antiparallel to each other. This places the maximum distance between electron density on adjacent shells in order to minimize electron repulsion.

As in the case of neon, if we consider the tetrahedral full-shell configuration, then argon will have a nested spherical tetrahedral orbital arrangement, the 2^nd and 3^rd shells each containing four di-electrons (below, right).

Intuitively, this seems more stable as it involves a greater degree of field cancellation than a full p-orbital with a single electron occupying each lobe. The entire shell contains electron density. It will be highest at the center of the face of each orbital (as in the traditional sp³ lobe shapes) and will decrease toward the nodal regions between orbitals, where electron density will be lowest (though not zero). Together, all atomic orbitals constitute a single, phase-locked, resonant, coherent, harmonic, stationary wave.

As mentioned above, the nested tetrahedral shells align in an opposite or antiparallel fashion. The reason this geometry is so stable for electrons is that the region of lowest electron density on one shell lines up with the region of maximum electron density on the adjacent shell, minimizing repulsion. In the diagram (below), the vertex represents where the nodes between orbitals intersect. This is a point of lowest electron density. It lies directly over the center of the face of the orbital beneath it, which is the point of highest electron density in that orbital. This is therefore the lowest energy state that nested tetrahedral shells can achieve.

---

RETURN to the Periodic Table or COMPARE to Neon.

SEE OTHER NOBLE GASES: Helium, Neon, Argon, Krypton, Xenon

RETURN to Periodic Table

Chlorine is the 17^th element on the periodic table. It has 17 protons and (mostly) 18 neutrons in the nucleus, giving it a mass of 35 amu, and it has 17 electrons enveloping the nucleus.

Electron Shell

Chlorine has the same electron configuration as fluorine, just one shell larger. This allows chlorine a broader palette of chemical reactions, given the presence of a d-orbital. Like sulfur, this gives rise to more possible molecular geometries. Chlorine can sustain multiply bonded atoms, but it takes bonding atoms with stronger electronegativity than chlorine in order to draw it into multiple bonds. This limits to list of contenders essentially to oxygen and fluorine, as we see in the chlorate (ClO₃^–) ion. (The wireframe indicates the boundary of the n=3 shell.)

NOTE: The small spheres in the image above simply indicate the directions of maximum electron density. The 3^rd shell orbitals themselves will be more like spherical tetrahedra. The entire shell will be filled with electron density. It will be highest at the center of the face of each orbital (as in the traditional sp³ lobe shapes) and will decrease toward the nodal regions between orbitals, where electron density will be lowest (though not zero). Like argon, chlorine features two nested spherical tetrahedra, the inner 2^nd shell comprising 4 di-electrons, the outer 3^rd shell comprising 3 di-electrons and 1 single electron. In this case the three orbitals containing the di-electrons (dark pink) will occupy slightly more volume than the one containing an unpaired electron.

Bonding & ion formation

Chlorine is keen to obtain an extra electron to fill its third shell and it can bond with many atoms on the periodic table. Chlorine can make one or more covalent bonds or gain an electron in ionic interaction in order to reach the stability of the 3s²3p⁶ noble gas configuration of argon, which is a multi-di-electron state with three concentric full shells. That is why chlorine forms a 1– ionic state. The negative ion is larger than its neutral version because electrons now outnumber protons. This results in a lower effective nuclear charge — a lower average attraction by the nucleus on each electron.

Salt

When sodium and chlorine interact, sodium gives the electron it wants to lose to chlorine, which is keen to gain it. This forms both atoms into their ions and allows both to achieve full shell configurations. The ions can then stick to each other because of their opposite charges, forming sodium chloride (NaCl) crystals. This process is called ionic bonding, and it occurs between a metal (from the left side of the periodic table) and a non-metal (from the right side). The term “salt” can also be used to apply to any ionic crystal.

Below is another visualization of the formation of sodium chloride from the point of view of electronegativity. The electrons in a bond will always be pulled more strongly towards the atom that is closer to the violet end of the electronegativity spectrum. In this case, there is such a great difference between how strongly they hold their electrons that the chlorine will be able to strip the sodium’s electron completely away. This clarifies why the two ions form, as opposed to them still sharing the electrons (as occurs in water).

Sodium chloride (NaCl) dissolves in water because the polar H₂O molecules and the ions in the crystal attract each other. The water molecules can therefore tug ions off the crystal and still satisfy the ion’s desire to attract their opposite polarity. As each ion leaves the crystal, it becomes hydrated — surrounded by water molecules.

When the water is allowed to evaporate from the salt solution, the ions become increasingly exposed to one another, and the solid crystals re-form due to electrostatic attraction.

Although the ocean contains many different ions and just about every element on the periodic table (in trace amounts), it is mainly made up of water (H₂O) and sodium chloride (NaCl). Of these four elements, the mass of sea water is made up of about 86% oxygen, 11% hydrogen, 1.9% chloride, and 1.1% sodium.

---

RETURN to the Periodic Table

OTHER GROUP VII HALOGENS: Fluorine, Chlorine, Bromine, Iodine

RETURN to Periodic Table

Sulfur is the 16^th element on the periodic table. It has 16 protons and 16 neutrons in the nucleus, giving it a mass of 32 amu, and it has 16 electrons enveloping the nucleus.

Electron Shell

Sulfur has the same electron configuration as oxygen, though as a larger atom and one with a d-orbital into which it can hybridize, sulfur has different properties and can make different molecular geometries than oxygen. (The wireframe indicates the boundary of the n=3 shell.)

NOTE: The small spheres in the image above simply indicate the directions of maximum electron density. The 3^rd shell orbitals themselves will be more like spherical tetrahedra. The entire shell will be filled with electron density. It will be highest at the center of the face of each orbital (as in the traditional sp³ lobe shapes) and will decrease toward the nodal regions between orbitals, where electron density will be lowest (though not zero). Like argon, sulfur features two nested spherical tetrahedra, the inner 2^nd shell comprising 4 di-electrons, the outer 3^rd shell comprising 2 di-electrons and 2 single electrons. In this case the two orbitals containing the di-electrons (dark pink) will occupy slightly more volume than the two containing an unpaired electron.

Bonding & ion formation

As the structure below indicates, it is common for sulfur to make two bonds. In its natural crystalline form, sulfur forms S₈ rings where each sulfur atoms is bonded to two adjacent sulfur atoms. Like oxygen (though not quite as strongly), sulfur is keen to gain two electrons in an ionic interaction in order to reach the stability of the 3s²3p⁶ noble gas configuration of argon, which is a multi-di-electron state with three concentric full shells. That is why sulfur forms a 2– ionic state. The negative ion is much larger than its neutral version because electrons now outnumber protons by two. This results in a lower effective nuclear charge — a lower average attraction by the nucleus on each electron, which expands the size of the electron shell as it is attracted inward with less force.

But if sulfur is approached by highly electronegative atoms like oxygen or fluorine, they can induce sulfur’s paired valence electrons to uncouple, yielding 6 unpaired electrons available for bonding. This is only possible because the 3^rd shell contains a d-orbital. While sulfur does not usually have electrons in its 3d-orbital, the 3^rd shell has enough volume to accommodate those orbitals. This allows sulfur to hybridize into those empty d-orbitals and to make up to 6 covalent bonds, as we see in the octahedral SF₆ molecule.

Uses

Sulfur is an abundant element and has a great many uses in nature, biology and chemical technology.

Sulfur is a relatively common atom in protein structures, and it is therefore present throughout the nutrient cycle, as well as being a waste product from the purification of natural gas and oil, which derive from organic materials.

One common form of sulfur in nature is in sulfide molecules, for example hydrogen sulfide (H₂S) gas, which is experienced as that notorious ‘rotten egg’ smell. Another is the sulfur dioxide (SO₂) gas.

One of the most stable forms of sulfur in nature is the sulfate ion (SO₄^2–). When two electrons are made available by a nearby metal atom, 1 sulfur atom and 4 oxygen atoms can form a symmetrical, tetrahedral arrangement where every atom has a full and symmetrical outer electron shell. Sulfates are a key nutrient in nature.

---

RETURN to the Periodic Table

RETURN to Periodic Table

Phosphorus is the 15^th element on the periodic table. It has 15 protons and 16 neutrons in the nucleus, giving it a mass of 31 amu, and it has 15 electrons enveloping the nucleus.

Electron Shell

Phosphorus has the same tetrahedral sp³ electron configuration as nitrogen, though, as a larger atom, the chemistry of phosphorus has important differences from that of nitrogen. With an available d-orbital into which to hybridize, phosphorus can achieve a greater variety of molecular geometries than can nitrogen. (The wireframe indicates the boundary of the n=3 shell.)

NOTE: The small spheres in the image above simply indicate the directions of maximum electron density. The 3^rd shell orbitals themselves will be more like spherical tetrahedra. The entire shell will be filled with electron density. It will be highest at the center of the face of each orbital (as in the traditional sp³ lobe shapes) and will decrease toward the nodal regions between orbitals, where electron density will be lowest (though not zero). Like argon, phosphorus features two nested spherical tetrahedra — an inner 2^nd shell comprising 4 di-electrons, and the outer 3^rd shell comprising 1 di-electron and 3 single electrons. In this case the orbital containing the di-electron (dark pink) will occupy slightly more volume than the three containing an unpaired electron.

Bonding & ion formation

Phosphorus is keen to obtain three extra electrons to fill its 3^rd shell, though it is not as pressing of an issue for phosphorus as it is for nitrogen since the nuclear charge is lower at the boundary of the larger atom. Phosphorus can make one or more covalent bonds or it can gain three electrons in ionic interaction in order to reach the stability of the 3s²3p⁶ noble gas configuration of argon, which is a multi-di-electron state with three concentric full shells. That is why phosphorus forms a 3– ionic state. The negative ion is much larger than its neutral version because electrons now outnumber protons by three. This results in a much lower effective nuclear charge — a lower average attraction by the nucleus on each electron, which expands the size of the electron shell as it is attracted inward with less force.

But if phosphorus is approached by highly electronegative atoms like oxygen or chlorine, they can induce phosphorus’s paired valence electrons to uncouple, yielding 5 unpaired electrons available for bonding. This is only possible because the 3^rd shell contains a d-orbital. While phosphorus does not usually have electrons in its 3d-orbital, the 3^rd shell has enough volume to accommodate those orbitals. This allows phosphorus to hybridize into that space and to make up to 5 covalent bonds, as we see in the trigonal bipyramidal PCl₅ molecule.

Uses

Since phosphorus has a lower electronegativity than nitrogen, it’s outer electrons are not as well bound. Oxygen will therefore react more easily and strongly with phosphorus than with nitrogen, which is why phosphorus burns (combusts) so readily. It can therefore be used to make matches and incendiary devices.

The most stable form of phosphorus in nature is the phosphate ion (PO₄^3–). When three electrons are made available by nearby metal atoms, 1 phosphorus atom and 4 oxygen atoms can form a symmetrical, tetrahedral arrangement where every atom has a full and symmetrical outer electron shell. Phosphates are a key nutrient in nature. They are commonly found in rocks, from where they are released by erosion, and they also help make up the ‘backbone’ of the DNA molecule in all living things.

---

RETURN to the Periodic Table

RETURN to Periodic Table

Silicon is the 14^th element on the periodic table. It has 14 protons and 14 neutrons in the nucleus, giving it a mass of 28 amu, and it has 14 electrons enveloping the nucleus.

Electron Shell

Silicon has the same electron structure as carbon, a hybridized, tetrahedral sp³ configuration. (The wireframe indicates the boundary of the n=3 shell.)

NOTE: The small spheres in the image above simply indicate the directions of maximum electron density. The 3^rd shell orbitals themselves will be more like spherical tetrahedra. The entire shell will be filled with electron density. It will be highest at the center of the face of each orbital (as in the traditional sp³ lobe shapes) and will decrease toward the nodal regions between orbitals, where electron density will be lowest (though not necessarily zero). Like argon, silicon features two nested spherical tetrahedra — an inner 2^nd shell comprising 4 di-electrons, and an outer 3^rd shell comprising 4 single electrons.

Bonding & Properties

Since silicon is a similar though larger atom than carbon, the four degenerate 3^rd shell electrons form a less coherent quantum system than they do in carbon. The larger size gives pure silicon different (more semi-metallic) properties than carbon. This is because metals are characterized by the ease with which their electrons can delocalize from the atomic core. The less well-bound they are to the nucleus, the more metallic the nature of the element. This means elements become more metallic as we move down the periodic table (as their ionization energies decrease).

Like carbon silicon can make 4 covalent bonds, and its natural crystal form has the same network covalent structure as diamond. This gives it a high melting point, but not the same strength or hardness as diamond because, being larger, silicon does not bond with itself as strongly as carbon does. We also do not see many organic molecule analogues involving silicon in place of carbon.

Uses

With a lower electronegativity than carbon, oxygen bonds more strongly to silicon than to carbon, making silicates (sand) one of the primary components of earth’s crust. Due to its larger size and less coherent hybridization, along with the presence of a d-orbital in the 3^rd shell into which to expand the electron octet, silicon is also capable of octahedral bonding. Silicon-based materials (like silicates) can therefore have more complex crystalline structures than carbon-based oxides.

Since silicon becomes less of an electrical resistor as temperature increases, it is a semiconductor, with important uses in electronics and photovoltaics. These properties are why silicon became the foundation of modern electronics. Silicon can be made to conduct electricity a little, a lot, or hardly at all, simply by adding tiny amounts of other elements, which makes it easy to control how electric current flows. Just as importantly, silicon naturally forms a very stable layer of silicon dioxide on its surface, which acts as an excellent electrical insulator and protective coating. This combination — precise control of electrical behavior, a reliable insulating layer, and the fact that silicon is abundant, durable, and well-understood—has made it ideal for manufacturing the billions of tiny switches (transistors) that power computers, phones, and other electronic devices.

---

RETURN to the Periodic Table

RETURN to Periodic Table

Aluminium (or Aluminum) is the 13^th element on the periodic table. It has 13 protons and 14 neutrons in the nucleus, giving it a mass of 27 amu, and it has 13 electrons enveloping the nucleus.

Electron Shell

Aluminium has interesting chemistry and bears similarities to several different elements. It lies in Group III beneath boron, and with a 3p¹ electron configuration, it is similar to boron’s 2p¹ configuration. This implies that it will form a triangular (sp² trigonal planar) geometry. It also has similarities to beryllium and can form bonds of a covalent nature. However, since aluminium is larger than boron, its valence electrons fill a larger volume and are therefore less coherent as a single quantum system. This makes aluminium more metallic in nature than the semi-metallic boron above it. (The wireframe indicates the boundary of the n=3 shell.)

NOTE: The small spheres in the image above simply indicate the directions of maximum electron density. The 3^rd shell orbitals themselves will be more like three longitudinal sections. The entire shell will be filled with electron density. It will be highest at the center of the face of each orbital (as in the traditional sp³ lobe shapes) and will decrease toward the nodal regions between orbitals, where electron density will be lowest (though not zero).

Each of these three hybrid orbitals contains one electron. Like boron’s configuration, this arrangement is symmetrical in the equatorial plane but does not have equivalent symmetry in all directions.

Bonding & ion formation

Aluminum will bond with other metal atoms via a metallic bond, and with non-metals via an ionic bond in which it loses its three valence electrons. It will lose all three at once in order to reach the stability of a full 2^nd shell. This is the same electron configuration as the 2s²2p⁶ noble gas configuration of neon — a multi-di-electron state with two concentric full shells. That is why aluminium forms a 3⁺ ionic state, as in aluminum oxide (Al₂O₃).

The positive ion is much smaller than the neutral atom because protons now outnumber electrons by three. This results in a much higher effective nuclear charge — a higher average attraction by the nucleus on each electron, which shrinks the size of the electron shell as it is attracted inward with more force.

Neutral aluminium (Al) atom (L) compared to the much smaller Al³⁺ ion (R)

Some uses & properties

Unlike most other small ions, aluminium ions do not play a biologically useful role. In fact, given their small size and high charge, Al³⁺ ions are toxic to most organisms.

Aluminium occurs in nature in a bonded, crystalline form. Reducing it to pure metal is done via electrolysis, and uses a large amount of electricity because it requires 3 electrons for every atom of aluminium to turn Al³⁺ ions into Al⁰ metal. One of the key uses of this light and strong metal is in the manufacture of aircraft. The first aircraft engine, built by the Wright brothers, was made out of aluminium specifically due to its low density, making for a lighter engine.

When we compare aluminium to boron (B), we see a significant size difference (see below). We also see a significant difference between the magnetic susceptibility of aluminium and the other Group III elements, since aluminium is paramagnetic (χ_m=+16.5) while the others are diamagnetic. Aluminium’s paramagnetic strength is also very similar to that of sodium (Na) (χ_m=+16).

Size comparison between boron (left) and aluminium (right)

This pattern might suggest that, in its crystalline metallic state, aluminium’s electronic structure and band-filling are somewhat unusual compared to a simple “three free electrons per atom” picture. In other words, while aluminium is conventionally treated as a trivalent metal in the band description, it may still be possible for certain bulk properties (especially transport and magnetic response) to be influenced disproportionately by one “more mobile” conduction channel, with the remaining valence character contributing differently through band structure, screening, and bonding. (This would not imply that two electrons remain literally localized on each atom, but rather that the metallic state can reorganize the valence character in ways that do not map one-to-one onto the isolated-atom picture. In compounds, by contrast, differences in electronegativity and bonding strongly favor aluminium’s familiar trivalent ionic behavior.)

If this speculation is even partially on the right track, it could help frame the following observations in a cautious way:

• aluminium has a magnetic susceptibility value similar to sodium (Na).

• aluminium has a 2^nd ionization energy whose increase is notably larger than the elements on either side of it (Mg & Si). While ionization energies are atomic (gas-phase) quantities and do not directly determine metallic conduction, the size of this jump is at least consistent with the idea that removing additional valence electrons from aluminium after the first can be comparatively “costly” in the isolated-atom sense, even if the metallic state distributes these electrons into bands rather than keeping them as discrete atomic electrons.

• aluminium is a soft metal. It seems reasonable that bond strength and hardness are shaped not only by electron count, but by how the bands fill, how effectively the electrons screen ion cores, and how easily the lattice accommodates dislocations. In that light, an “effective” carrier picture in which one portion of the valence response dominates conduction does not contradict aluminium’s softness, and may even be compatible with the idea that aluminium’s metallic bonding is comparatively less stiff than one might naively expect from a purely trivalent caricature.

• If aluminium behaves, in some measurements, as though a smaller subset of carriers dominates electrical transport — somewhat reminiscent (in effect, not in literal electron count) of copper (Cu), silver (Ag), and gold (Au) — that could help motivate why aluminium is among the very best electrical conductors in practice, typically ranking just after silver, copper, and gold, despite being a much lighter and more reactive metal. (Here, the main point is not that aluminium truly has “only one delocalized electron per atom,” but that its conduction and magnetic signatures can be shaped by band structure in a way that sometimes makes it look more “one-carrier-like” than a naive valence tally would suggest.)

Boron’s diamagnetism (χ_m=–6.7), on the other hand, would appear to emerge from the fact that it usually makes covalent bonds, in which all of its electrons are paired

---

RETURN to the Periodic Table

RETURN to Periodic Table

Magnesium is the 12^th element on the periodic table. It has 12 protons and 12 neutrons in the nucleus, giving it a mass of 24 amu, and it has 12 electrons enveloping the nucleus.

Electron Shell

Magnesium has a di-electron making up its 3^rd shell, with two full core shells within that have the identical configuration to neon. Given its larger size, the 3s² di-electron is not as well-bound as the 2s² di-electron on beryllium. Magnesium will therefore donate its valence electrons more readily than beryllium, but it will not be as reactive in doing so as the metals below it in Group II such as calcium, strontium, or barium.

As we saw in the case of neon, the 2^nd shell orbitals are more like spherical tetrahedra, and the 3^rd shell is a di-electron in a spherical s-orbital. These orbitals represent phase-locked, resonant, coherent, harmonic, stationary waves.

Ion formation

Magnesium is more willing to lose its two valence electrons in an ionic interaction than is beryllium because it is larger, and it will lose both at the same time in order to reach the stability of a full 2^nd shell. This is the same electron configuration as the 2s²2p⁶ noble gas configuration of neon — a multi-di-electron state with two concentric full shells. That is why magnesium forms a 2+ ionic state.

---

RETURN to the Periodic Table

OTHER GROUP II ELEMENTS: Beryllium, Magnesium, Calcium, Strontium, Barium